Covalent Compounds

1. Covalent Compounds

2. Covalent compounds contain covalent bonds • Covalent bonds = sharing electrons • Covalent bonds usually form between nonmetals. • Covalent bonds can involve multiple pairs of electrons: single, double, triple bonds.

3. Properties of covalent compounds: • Covalent compounds have low melting and boiling points. • Covalent compounds are usually soft, not brittle. • Covalent compounds are poor conductors.

4. Covalent Bonding Covalent bonds form by sharing electrons between atoms … … so that each atom appears to have an octet of electrons. Diatomic elements are good examples of covalent bonding.

5. The Diatomic Elements are: H2 N2 O2 F2 Cl2 Br2 I2 Known as the “hairogens”: H, N & O, halogens air H ogens air N and O = ??? …hence, the

6. Bonding in the Halogens F + F F2 F F Formation of a F-F bond

7. Bonding in the Halogens F + F F2 F F The overlap of two p-orbitals creates the single bond between fluorine atoms. F - F

8. See how a double bond occurs in an oxygen molecule and a triple bond occurs in a nitrogen molecule.

9. Bonding in Oxygen O + O O2 O O The overlap of four p-orbitals creates the double bond between oxygen atoms. O = O

10. N N Bonding in Nitrogen N + N N2 N N The overlap of six p-orbitals creates the triple bond between nitrogen atoms.

11. Comparison of single, double and triple bonds: Bond length: s--i--n--g--l--e > d o u b l e > triple A BA BA B Bond strength: single < double < triple

12. Covalent bonds result from the overlap of orbitals.

13. Covalent bonds result from the overlap of orbitals.

14. x z Recall the shapes of the three p-orbitals

15. Consider two p-orbitals from two different elements: As the orbital get closer …

16. Consider two p-orbitals from two different elements: A bond occurs when the orbitals overlap end to end.

17. Gilbert N. Lewis • American chemist and educator. • Defined acids as electron pair acceptors and bases as electron pair donors. • Explained his theory with “electron dot diagrams”. • Still in use today to explain molecular structure as well as acids and bases.

18. Writing Lewis Structures • Add up all of the valence electrons • Decide on a central atom. It has the lowest EN. H is never a central atom; halogens rarely are. • Draw the skeleton of the molecule and connect each symbol with a dash to indicate a bonding pair of electrons

19. Writing Lewis Structures • Complete the octet of the terminal atoms, add all the electrons and compare to #1 • Add any additional electrons to the central atom, even if it means having more than 8. • If there are not enough electrons to give every element an octet, consider multiple bonds.

20. Writing Lewis Structures Some things to remember: • Hydrogen can only have two electrons around it, not an octet. • The central atom is frequently the one that there is only 1 of. • Halogens are almost never the central atom and they never have double or triple bonds!

21. Write the Lewis structures for the following compounds: H2O CH4 OF2 PCl3 HCN CO CO2 SCl4 PCl5 XeCl4

22. Polyatomic Ions • An ion with two or more atoms. • Polyatomic ions have unique formulas and names: OH- = hydroxide ion SO42- = sulfate ion PO43- = phosphate ion

23. Lewis Structures of polyatomic ions • Write the Lewis structure just as you would for a compound, except … • …the number of valence electrons must be increased (or decreased) because of the charge on the ion. 5+4(6)+3 = 32 electrons PO43- Consider the phosphate ion. It has three extra electrons.

24. Draw the Lewis structure of the following polyatomic ions: • Hydroxide ion OH - • Sulfate ion SO42- • Phosphate ion PO43- • Nitrate ion NO31- • Ammonium ion NH4+